1.5 Oxidation and - PowerPoint Presentation

Slide1

1.5 Oxidation and

ReductionSlide2

Learning Outcomes

Introduction to oxidation

and reduction

: simple examples

only, e.g

. Na with Cl

2

, Mg with

O

2

, Zn

with Cu

2+.

Oxidation

and reduction

in terms of loss

and gain

of electrons.

Oxidising and reducing agents.

The electrochemical series as

a series

of metals arranged

in order

of their ability to

be oxidised

(reactions, other

than displacement

reactions,

not required

).

Electrolysis of (

i

) copper

sulfate

solution

with copper

electrodes and

(ii) acidified water

with inert

electrodes. (Half

equations only

required.)Slide3

Oxidation and reduction

Oxidation = addition of oxygen to a substance

C + O

2



CO

2

Reduction is loss of oxygen or addition of hydrogen

CuO + H

2

 Cu + H

2

OSlide4

examples

Sodium + chlorine

 sodium chloride

Na +

Cl

 Na

+

+

Cl

-

Na loses an electron [oxidised

]

Cl

gains an electron [reduced]Slide5
Slide6

Example 2

Magnesium +oxygen

 magnesium oxide

Mg + O 

MgO

Mg  Mg

+2

loses 2 electrons [oxidation]

O  O

-2

gains 2 electrons [reduction]Slide7

Example 3

Zinc +copper sulphate



Zinc sulphate+ Copper

Zn + Cu

+2

 Zn

+2

+ Cu

Zinc loses electrons (oxidised)

Copper gains electrons (reduced)Slide8

Oxidising agent

A substance that causes oxidation in another substanceSlide9

Reducing agent

A substance that causes reduction in another substance.Slide10

Oxidation is

loss of electrons;

R

eduction

is gain of electrons

CuO + H

2

 Cu + H

2

OCuO  Cu+2 and O-2Cu+2  Cu [gains 2 electrons] reduced

H

2

 H

2

+2

[loses 2 electrons] oxidised

O

-2

 O

-2 [ no change]Slide11

Oxidation numbers

The charge that an atom has or appears to have assuming that the compound is ionic.

Electrons always go the the most electronegative elementSlide12

Oxidation number rules 1

Elements on

their

own

=

0

H

2

= 0

Zn = 0

Cl2 = 0Slide13

Oxidation number rules 2

Ions = same as charge

Cu

+2

= +2

O

-2

= -2

Cl

-1

= -1Slide14

Oxidation number rules 3

Charges of all elements in a compound = 0

CuSO

4

Cu = +2

S = +6

O

4

= -8 [O = -2]

Total = +2 +6 –8 = 0Slide15

Oxidation number rules 4

Oxygen

= -2

Exceptions are

peroxides O = -1 [H

2

O

2

, Na

2

O2 ]OF2

O = +2, F = -1Slide16

Oxidation number rules 5

Hydrogen = +1

Exceptions are the metal hydrides

NaH

Na

= +1,

H

= -1Slide17

Oxidation number rules 6

Halogens [ Cl, F, I, Br] are always –1 except when joined to more electropositice element

Cl

2

O

Cl

= +1, O = -2Slide18

Oxidation number rules 7

In a complex ion the sum of all the charges = the chartge on the ion.

SO

4

-

2

S = +6, O

4

= -8 [O = -2]

+6 –8 = -2Slide19

redox

Oxidation is an increase on oxidation number

Reduction is a decrease in oxidation number.Slide20

Electrochemical Series

Electrochemical Series – Elements listed in order of ability to be

oxidisedSlide21

Metals

King [K]

Neptune [Na]

Caught [Ca]

Many [Mg]

Angry [Al]

Zulus [Zn]

Fighting [Fe]

Police [

Pb

] Constables [Cu] Having [Hg] Asthma [Ag] Attacks [Au]Slide22

Metals above hydrogen in the Reactivity Series react with acids to produce hydrogen gas

.

Zinc

Potassium

SodiumSlide23

Displacement of metals

Displacement reactions occur when a metal from the electrochemical series is mixed with the ions of a metal lower down in the series. The

atoms

of the more reactive metal push their electrons on to

ions

of the less reactive metal.Slide24

Displacement

More reactive metal displaces less reactive from a solution

Mg + CuSO

4

= MgSO

4

+ Cu

Mg + Cu

+2

 Mg+2 + CuMg loses electrons (Oxidised)Cu+2 gains electrons (reduced)Slide25

Learning Outcomes

Rusting of iron.

Swimming-pool water treatment.

Use of scrap iron to

extract copper

.

Electroplating. Purification

of copper.

Chrome and nickel

plating. Cutlery

.Slide26

Rust

Rust

is the formation of iron oxides (usually red oxides), formed by the reaction of iron and oxygen in the presence of water or air moisture.

OxidationSlide27

Swimming pools

The water in swimming pools is kept sterile by the addition of oxidizing agents, chlorine or chlorine compounds, which kill microorganisms by oxidation. The active agent is usually

chloric

(1) acid (

HOCl

). It may be formed in two ways

1. Direct chlorination of the water:

Cl

2(aq)

+H

2O (l)  HOCl

(aq)

+ Cl

−

(aq)

+ H

+

(aq)

Note that when the Cl

2

reacts with the water it is both oxidized and reducedSlide28

Swimming pools

2. The addition of sodium chlorate(I) [sodium hypochlorite]:

NaOCl

(s)

+ H

2

O

(l)

Na

+

(aq) + OH− (aq) + HOCl (aq) Nowadays chlorine is not used, mainly on grounds of safety. Pools are sterilized with chlorine compounds, which produce

chloric

(I) acid when they dissolve in water. These compounds act in essentially the same way as chlorine. Sodium chlorate(I) is one such compound.Slide29

Use of scrap iron to extract copper.

(Dissolved CuSO

4

) + (Metallic Fe) ==> (Dissolved FeSO

4

) + (Metallic Cu)Slide30

Electrolysis

Chemical reaction caused by the passage of an electric current through a liquid known as the electrolyteSlide31

Definitions

Electrolyte -

liquid in which electrolysis takes place. Usually an ionic solution but it can also be a fused [melted] ionic compound

Anode -

positive electrode. Positive because the battery sucks electrons out of it

Cathode.

Negative electrode. Negative because the battery pumps electrons into it.

Anion - negative ion.

Called anion because it is attracted to the opposite charge of the anode

Cation

- positive ion.

Called

cation

because it is attracted to the opposite charge of the cathode.

Inert Electrodes -

do not react with the electrolyte Graphite and Pt

Active electrodes -

react with electrolyte e.g. Copper and iron

 Slide32

ElectrolysisSlide33

Electroplating

Electroplating

Covering cathode in metal e.g. Cu by making it cathode in copper sulphate solutionSlide34

Copper platingSlide35

Copper Plating

Anode reaction

Cu

(s)

= Cu

2+

(

aq

)

+ 2e

-Anode loses mass as copper dissolves offImpurities [Au, Ag, Pt etc.] fall to bottomCathode reaction

Cu

2+

(aq)

+ 2e

-

= Cu

(s)

Cathode gains mass as Cu is deposited on it

Cu is 99.9% pure

 Slide36

Learning Outcomes

Mandatory experiment

1.2 (half

equations only

required, e.g

. 2Br

–

– 2e

–

→ Br

2

).

Demonstration of

ionic movement

.

Demonstration of electrolysis

of aqueous

sodium

sulfate

(

using universal

indicator)

and of aqueous

potassium iodide (

using phenolphthalein

indicator)

with inert

electrodes. (Half

equations only

required.)Slide37

Ionic Movement

During electrolysis of a solution of Copper Chromate in dil. Hydrochloric acid, positive ions (

cations

) are attracted to the negative electrode (cathode) and negative ions (anions) are attracted to the positive electrode (anode). If these ions are coloured, their movement may be observed visually.

Examples of coloured ions include;

copper(II) [Cu

2+

] - blue

chromate(VI) [CrO

42-

] – yellow Slide38

Q & A to Ionic Movement

Expt

(1) What colour is the copper(II) chromate solution?

Copper(II) chromate solution is an olive green colour.

(2) What colour is observed at the positive electrode after the power supply has been turned on for some time?

A yellow colour is observed at the positive electrode.

(3) What colour is observed at the negative electrode after the power supply has been turned on for some time?

A blue colour is observed at the negative electrode.

(4) Explain in terms of the movement of ions why different colours are formed at each electrode.

When the circuit is completed, positive copper ions (Cu

2+) are attracted to the negative electrode. These ions have a blue colour. Similarly negative chromate(VI) ions (CrO

42-

) are attracted to the positive electrode. These ions are coloured orange.

(5) What is the function of the dilute hydrochloric acid?

The dilute hydrochloric acid is required to complete the circuit. Slide39

Electrolysis of Sodium sulphate

S

olution

of Na

2

SO

4

+

universal

indicator

H

+

ions are produced at the positive electrode (oxidation of O

2-

in water) while OH

-

ions are produced at the negative electrode as the H

+

in water is reduced to H

2

(g).Slide40

Sodium Sulphate and Universal IndicatorSlide41

Electrolysis of Sodium Sulphate

Red is acid at the positive

electrode

2H

2

O(l)



O

2

(g) + 4H

+

(

aq

) + 4 e

-

lose electrons = oxidation = anode

Purple is base

at

the negative

electrode

H

2

O(l) + 2 e

-



H

2

(g) + 2OH

-

(

aq

)

gain electrons = reduction = cathodeSlide42

Electrolysis of

P

otassium Iodide

Solution of

KI

+ phenolphthalein

Brown

I

2

(s) forms

at

the positive electrode and some yellow/orange I

3

-

forms in solution. At the negative electrode, H

+

is again reduced to H

2

(g) and the phenolphthalein turns pink due to the OH

-

ions.Slide43

Electrolysis of Potassium IodideSlide44

Electrolysis of Potassium Iodide

KI

 K

+

+ I-

Iodide loses electrons  Brown iodine

2I-  I

2

+ 2e

-

Anode, OxidationH

2

O  H

+

+ OH

-

OH

-

is basic , Phenolphthalein Purple.